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Calorimetry

Calorimetry

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Calorimetry

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  1. Calorimetry

  2. D(PE) (Products) Burning of a Match System Surroundings (Reactants) Potential energy Energy released to the surrounding as heat Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 293

  3. Endothermic Reaction Reactant + Energy Product Surroundings System Conservation of Energy in a Chemical Reaction In this example, the energy of the reactants and products increases, while the energy of the surroundings decreases. In every case, however, the total energy does not change. Surroundings Energy System Before reaction After reaction Myers, Oldham, Tocci, Chemistry, 2004, page 41

  4. Exothermic Reaction Reactant Product + Energy Surroundings System Conservation of Energy in a Chemical Reaction In this example, the energy of the reactants and products decreases, while the energy of the surroundings increases. In every case, however, the total energy does not change. Surroundings System Energy Before reaction After reaction Myers, Oldham, Tocci, Chemistry, 2004, page 41

  5. H2O(s) + heat  H2O(l) H2O(l)  H2O(s) + heat melting freezing System Direction of Heat Flow Surroundings EXOthermic qsys < 0 ENDOthermic qsys > 0 System Kotz, Purcell, Chemistry & Chemical Reactivity 1991, page 207

  6. Caloric Values Food joules/grams calories/gram Calories/gram Protein 17 000 4000 4 Fat 38 000 9000 9 Carbohydrates 17 000 4000 4 1calories = 4.184 joules 1000 calories = 1 Calorie "science" "food" Smoot, Smith, Price, Chemistry A Modern Course, 1990, page 51

  7. Experimental Determination of Specific Heat of a Metal Typical apparatus used in this activity include a boiler (such as large glass beaker), a heat source (Bunsen burner or hot plate), a stand or tripod for the boiler, a calorimeter, thermometers, samples (typically samples of copper, aluminum, zinc, tin, or lead), tongs (or forceps or string) to handle samples, and a balance.

  8. Thermometer Styrofoam cover Styrofoam cups Stirrer A Coffee Cup Calorimeter Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 302

  9. Bomb Calorimeter thermometer stirrer full of water ignition wire steel “bomb” sample

  10. oxygen supply thermometer ignition wires stirrer magnifying eyepiece insulating jacket air space bucket heater crucible water ignition coil sample steel bomb 1997 Encyclopedia Britanica, Inc.

  11. A Bomb Calorimeter

  12. Causes of Change - Calorimetry Outline Keys http://www.unit5.org/chemistry/Matter.html

  13. Gas - KE Boiling - PE  Liquid - KE  Melting - PE  Solid - KE  Heating Curves 140 120 100 80 60 40 Temperature (oC) 20 0 -20 -40 -60 -80 -100 Time

  14. Heating Curves • Temperature Change • change in KE (molecular motion) • depends on heat capacity • Heat Capacity • energy required to raise the temp of 1 gram of a substance by 1°C • “Volcano” clip - • water has a very high heat capacity Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

  15. Heating Curves • Phase Change • change in PE (molecular arrangement) • temp remains constant • Heat of Fusion (Hfus) • energy required to melt 1 gram of a substance at its m.p. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

  16. Heating Curves • Heat of Vaporization (Hvap) • energy required to boil 1 gram of a substance at its b.p. • usually larger than Hfus…why? • EX: sweating, steam burns, the drinking bird Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

  17. Phase Diagrams • Show the phases of a substance at different temps and pressures. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

  18. Humor A small piece of ice which lived in a test tube fell in love with a Bunsen burner. “Bunsen! My flame! I melt whenever I see you” said the ice. The Bunsen burner replied” “It’s just a phase you’re going through”.

  19. Heating Curve for Water(Phase Diagram) F 140 120 100 80 60 40 20 0 -20 -40 -60 -80 -100 Heat= m x Cvap Cv = 2256 J/g E D BP Heat= m x Cfus Cf = 333 J/g Heat= m x DT x Cp, gas Cp (steam) = 2.042 J/goC Heat= m x DT x Cp, liquid Temperature (oC) Cp = 4.184 J/goC B MP C Heat= m x DT x Cp, solid A  B warm ice B  C melt ice (solid  liquid) C  D warm water D  E boil water (liquid  gas) E  D condense steam (gas  liquid) E  F superheat steam Cp (ice) = 2.077 J/goC A Heat

  20. Calculating Energy Changes - Heating Curve for Water 140 DH = mol xDHvap DH = mol xDHfus 120 100 80 Heat = mass xDt x Cp, gas 60 40 Temperature (oC) 20 Heat = mass xDt x Cp, liquid 0 -20 -40 -60 Heat = mass xDt x Cp, solid -80 -100 Time

  21. Equal Masses of Hot and Cold Water Thin metal wall Insulated box Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 291

  22. Water Molecules in Hot and Cold Water Hot water Cold Water 90 oC 10 oC Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 291

  23. Water Molecules in the same temperature water Water (50 oC) Water (50 oC) Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 291

  24. Heat Transfer Surroundings Final Temperature Block “B” Block “A” SYSTEM 20 g (20oC) 30oC 20 g (40oC) Al Al m = 20 g T = 40oC m = 20 g T = 20oC What will be the final temperature of the system ? a) 60oC b) 30oC c) 20oC d) ? Assume NO heat energy is “lost” to the surroundings from the system.

  25. Heat Transfer ? Surroundings Final Temperature Block “B” Block “A” SYSTEM 20 g (20oC) 30.0oC 20 g (40oC) 10 g (20oC) 33.3oC 20 g (40oC) Al Al m = 20 g T = 40oC m = 10 g T = 20oC What will be the final temperature of the system ? a) 60oC b) 30oC c) 20oC d) ? Assume NO heat energy is “lost” to the surroundings from the system.

  26. Heat Transfer Surroundings Final Temperature Block “B” Block “A” SYSTEM 20 g (20oC) 30.0oC 20 g (40oC) 10 g (20oC) 33.3oC 20 g (40oC) 10 g (40oC) 26.7oC 20 g (20oC) Al Al m = 20 g T = 20oC m = 10 g T = 40oC Assume NO heat energy is “lost” to the surroundings from the system.

  27. Heat Transfer Surroundings Final Temperature Block “B” Block “A” SYSTEM 20 g (20oC) 30.0oC 20 g (40oC) 10 g (20oC) 33.3oC 20 g (40oC) 10 g (40oC) 26.7oC 20 g (20oC) H2O Ag m = 75 g T = 25oC m = 30 g T = 100oC Real Final Temperature = 26.6oC Why? We’ve been assuming ALL materials transfer heat equally well.

  28. Specific Heat • Water and silver do not transfer heat equally well. Water has a specific heat Cp = 4.184 J/goC Silver has a specific heat Cp = 0.235 J/goC • What does that mean? It requires 4.184 Joules of energy to heat 1 gram of water 1oC and only 0.235 Joules of energy to heat 1 gram of silver 1oC. • Law of Conservation of Energy… In our situation (silver is “hot” and water is “cold”)… this means water heats up slowly and requires a lot of energy whereas silver will cool off quickly and not release much energy. • Lets look at the math!

  29. Specific Heat The amount of heat required to raise the temperature of one gram of substance by one degree Celsius.

  30. Calculations involving Specific Heat OR cp = Specific Heat q = Heat lost or gained T = Temperature change m = Mass

  31. Table of Specific Heats Specific Heats of Some Common Substances at 298.15 K Substance Specific heat J/(g.K) Water (l) 4.18 Water (s) 2.06 Water (g) 1.87 Ammonia (g) 2.09 Benzene (l) 1.74 Ethanol (l) 2.44 Ethanol (g) 1.42 Aluminum (s) 0.897 Calcium (s) 0.647 Carbon, graphite (s) 0.709 Copper (s) 0.385 Gold (s) 0.129 Iron (s) 0.449 Mercury (l) 0.140 Lead (s) 0.129

  32. Latent Heat of Phase Change Molar Heat of Fusion The energy that must be absorbed in order to convert one mole of solid to liquid at its melting point. The energy that must be removed in order to convert one mole of liquid to solid at its freezing point.

  33. Latent Heat of Phase Change #2 Molar Heat of Vaporization The energy that must be absorbed in order to convert one mole of liquid to gas at its boiling point. The energy that must be removed in order to convert one mole of gas to liquid at its condensation point.

  34. Latent Heat – Sample Problem Problem: The molar heat of fusion of water is 6.009 kJ/mol. How much energy is needed to convert 60 grams of ice at 0C to liquid water at 0C? Mass of ice Molar Mass of water Heat of fusion

  35. Heat of Reaction The amount of heat released or absorbed during a chemical reaction. Endothermic: Reactions in which energy is absorbed as the reaction proceeds. Exothermic: Reactions in which energy is released as the reaction proceeds.

  36. “loses” heat Calorimetry Surroundings SYSTEM Tfinal = 26.6oC H2O Ag m = 75 g T = 25oC m = 30 g T = 100oC

  37. Calorimetry Surroundings SYSTEM H2O Ag m = 75 g T = 25oC m = 30 g T = 100oC

  38. 1 BTU (British Thermal Unit) – amount of heat needed to raise 1 pound of water 1oF. 1 calorie - amount of heat needed to raise 1 gram of water 1oC 1 Calorie = 1000 calories “food” = “science” Candy bar 300 Calories = 300,000 calories English Joules 1 calorie = 4.184 Joules Metric = _______

  39. 140 DH = mol xDHvap DH = mol xDHfus 120 100 80 60 Heat = mass xDt x Cp, gas 40 Temperature (oC) 20 Heat = mass xDt x Cp, liquid 0 -20 -40 -60 Heat = mass xDt x Cp, solid -80 -100 Time Cp(ice) = 2.077 J/g oC It takes 2.077 Joules to raise 1 gram ice 1oC. X Joules to raise 10 gram ice 1oC. (10 g)(2.077 J/g oC) = 20.77 Joules X Joules to raise 10 gram ice 10oC. (10oC)(10 g)(2.077 J/g oC) = 207.7 Joules q = Cp . m .DT Heat = (specific heat) (mass) (change in temperature)

  40. 140 DH = mol xDHvap DH = mol xDHfus 120 100 80 60 Heat = mass xDt x Cp, gas 40 Temperature (oC) 20 Heat = mass xDt x Cp, liquid 0 -20 -40 -60 Heat = mass xDt x Cp, solid -80 -100 Time Given Ti= -30oC Tf = -20oC q = Cp . m .DT Heat = (specific heat) (mass) (change in temperature) q = 207.7 Joules

  41. T = 500oC Fe T = 20oC mass = 240 g 240 g of water (initially at 20oC) are mixed with an unknown mass of iron (initially at 500oC). When thermal equilibrium is reached, the system has a temperature of 42oC. Find the mass of the iron. mass = ? grams - LOSE heat = GAIN heat - [(Cp,Fe) (mass) (DT)] = (Cp,H2O) (mass) (DT) - [(0.4495 J/goC) (X g) (42oC - 500oC)] = (4.184 J/goC) (240 g) (42oC - 20oC)] - [(0.4495) (X) (-458)] = (4.184) (240 g) (22) Drop Units: 205.9 X = 22091 X = 107.3 g Fe Calorimetry Problems 2 question #5

  42. T = 785oC mass = 97 g Au T = 15oC mass = 323 g - LOSE heat = GAIN heat A 97 g sample of gold at 785oC is dropped into 323 g of water, which has an initial temperature of 15oC. If gold has a specific heat of 0.129 J/goC, what is the final temperature of the mixture? Assume that the gold experiences no change in state of matter. - [(Cp,Au) (mass) (DT)] = (Cp,H2O) (mass) (DT) - [(0.129 J/goC) (97 g) (Tf - 785oC)] = (4.184 J/goC) (323 g) (Tf - 15oC)] Drop Units: - [(12.5) (Tf - 785oC)] = (1.35 x 103) (Tf - 15oC)] -12.5 Tf + 9.82 x 103 = 1.35 x 103 Tf - 2.02 x 104 3 x 104 = 1.36 x 103 Tf Tf = 22.1oC Calorimetry Problems 2 question #8

  43. T = 72oC T = 13oC mass = 87 g mass = 59 g - LOSE heat = GAIN heat If 59 g of water at 13oC are mixed with 87 g of water at 72oC, find the final temperature of the system. - [(Cp,H2O) (mass) (DT)] = (Cp,H2O) (mass) (DT) - [(4.184 J/goC) (87 g) (Tf - 72oC)] = (4.184 J/goC) (59 g) (Tf - 13oC) Drop Units: - [(364.0) (Tf - 72oC)] = (246.8) (Tf - 13oC) -364 Tf + 26208 = 246.8 Tf - 3208 29416 = 610.8 Tf Tf = 48.2oC Calorimetry Problems 2 question #9

  44. 140 DH = mol xDHvap DH = mol xDHfus 120 100 80 60 Heat = mass xDt x Cp, gas 40 Temperature (oC) 20 Heat = mass xDt x Cp, liquid 0 -20 -40 T = -11oC -60 mass = 38 g Heat = mass xDt x Cp, solid ice -80 -100 Time T = 56oC mass = 214 g - LOSE heat = GAIN heat A 38 g sample of ice at -11oC is placed into 214 g of water at 56oC. Find the system's final temperature. D water cools B warm water C A melt ice warm ice D A C B - [(Cp,H2O(l)) (mass) (DT)] = (Cp,H2O(s)) (mass) (DT) + (Cf) (mass) + (Cp,H2O(l)) (mass) (DT) - [(4.184 J/goC)(214 g)(Tf - 56oC)] = (2.077 J/goC)(38 g)(11oC) + (333 J/g)(38 g) + (4.184 J/goC)(38 g)(Tf - 0oC) - [(895) (Tf - 56oC)] = 868 + 12654 + (159) (Tf)] - 895 Tf + 50141 = 868 + 12654 + 159 Tf - 895 Tf + 50141 = 13522 + 159 Tf 36619 = 1054 Tf Tf = 34.7oC Calorimetry Problems 2 question #10

  45. 140 DH = mol xDHvap DH = mol xDHfus 120 100 80 60 Heat = mass xDt x Cp, gas 40 Temperature (oC) 20 Heat = mass xDt x Cp, liquid 0 -20 -40 -60 Heat = mass xDt x Cp, solid -80 -100 Time 1102 1102 (1000 g = 1 kg) 238.4 g 25 g of 116oC steam are bubbled into 0.2384 kg of water at 8oC. Find the final temperature of the system. - [qA + qB + qC] = qD - [(Cp,H2O) (mass) (DT)] + (Cv,H2O) (mass) + (Cp,H2O) (mass) (DT) = [(Cp,H2O) (mass) (DT)] qD = (4.184 J/goC) (238.4 g) (Tf - 8oC) qD = - 997Tf - 7972 qA = [(Cp,H2O) (mass) (DT)] qC = [(Cp,H2O) (mass) (DT)] qB = (Cv,H2O) (mass) qA = [(2.042 J/goC) (25 g) (100o - 116oC)] qC = [(4.184 J/goC) (25 g) (Tf - 100oC)] qA = (2256 J/g) (25 g) qA = - 816.8 J qA = 104.5Tf - 10450 qA = - 56400 J - [qA + qB + qC] = qD - [ - 816.8 - 56400 + 104.5Tf - 10450] = 997Tf - 7972 816.8 + 56400 - 104.5Tf + 10450 = 997Tf - 7972 67667 - 104.5Tf = 997Tf - 7979 A 75646 = 1102Tf C B Tf = 68.6oC D Calorimetry Problems 2 question #11

  46. Ti = 25oC mass = 264 g Pb A 322 g sample of lead (specific heat = 0.138 J/goC) is placed into 264 g of water at 25oC. If the system's final temperature is 46oC, what was the initial temperature of the lead? T = ? oC mass = 322 g Pb Tf = 46oC - LOSE heat = GAIN heat - [(Cp,Pb) (mass) (DT)] = (Cp,H2O) (mass) (DT) - [(0.138 J/goC) (322 g) (46oC - Ti)] = (4.184 J/goC) (264 g) (46oC- 25oC)] Drop Units: - [(44.44) (46oC - Ti)] = (1104.6) (21oC)] - 2044 + 44.44 Ti = 23197 44.44 Ti = 25241 Ti = 568oC Calorimetry Problems 2 question #12

  47. Ti = 85oC mass = 68 g 458.2 458.2 A sample of ice at –12oC is placed into 68 g of water at 85oC. If the final temperature of the system is 24oC, what was the mass of the ice? T = -12oC mass = ? g H2O ice Tf = 24oC GAIN heat = - LOSE heat qA = [(Cp,H2O) (mass) (DT)] 24.9 m qA = [(2.077 J/goC) (mass) (12oC)] [ qA + qB + qC ] = - [(Cp,H2O) (mass) (DT)] [ qA + qB + qC ] = - [(4.184 J/goC) (68 g) (-61oC)] qB = (Cf,H2O) (mass) 333 m qB = (333 J/g) (mass) 458.2 m = - 17339 qC = [(Cp,H2O) (mass) (DT)] m = 37.8 g qC = [(4.184 J/goC) (mass) (24oC)] 100.3 m qTotal = qA + qB + qC 458.2 m Calorimetry Problems 2 question #13

  48. Activation Energy Endothermic Reaction Energy + Reactants  Products Products Energy +DH Endothermic Reactants Reaction progress

  49. Calorimetry Problems 1 Calorimetry 1 Calorimetry 1 Keys http://www.unit5.org/chemistry/Matter.html

  50. Calorimetry Problems 2 Calorimetry 2 Specific Heat Values Calorimetry 2 Specific Heat Values Keys http://www.unit5.org/chemistry/Matter.html